Biological Buffers • AppliChem

i n t r o d u c
Introduction
The buffer concept
Many biochemical processes are markedly impaired by even small changes in the concentrations of free H + ions. It is
therefore usually necessary to stabilise the H + concentration in vitro by adding a suitable buffer to the medium, without,
however, affecting the functioning of the system under investigation. A buffer keeps the pH of a solution constant by taking
up protons that are released during reactions, or by releasing protons when they are consumed by reactions. The observation
that partially neutralised solutions of weak acids or bases are resistant to changes in pH when small amounts of strong acids
or bases are added led to the concept of the ‘buffer’.
Buffer capacity
Buffers consist of an acid and its conjugated base. The quality of a buffer is determined by its buffer capacity, i.e. its resistance
to changes in pH when strong acids or bases are added. In other words: the buffer capacity corresponds to the amount
of H + or OH – ions that can be neutralised by the buffer. The buffer capacity is related to the buffer concentration. The
graph described by the relation of the pH to the addition of H + /OH – ions is called the titration curve. The point of in flection
of the curve corresponds to the pKa value. At this point, the buffer capacity is at its maximum at the pKa value. This point
therefore corresponds to the mid-point of the pH range covered by the buffer and is where the concentration of acid and
base is the same. In the area of this pH range, therefore, relatively large amounts of H + /OH – ions result in only small changes
in pH.
A basic principle is that a buffer that has a pH value of one pH unit above or below the pKa value loses so much buffer capacity
that it no longer has any real buffer function. Based on the Henderson-Hasselbalch equation
pH = pKa + log [A – ]/[HA]
for the calculation of the pH of a weak acid or alkaline solution, the ions in the water must also be taken into account when
working in pH ranges below 3.0 and above 11.0. Most biochemical reactions, however, take place in the pH range between
6.0 and 10.0.
The pH value
Conductivity can also be detected in highly purified water due to the OH – and H3O + ions resulting from the autoprotolysis of
water. This intrinsic dissociation of the water is an equilibrium reaction and the product of the concentrations of the two
ions represents a constant:
K = [H3O + ] x [OH – ].
This value is depends on the temperature only and is 10 –14 for purified water at 22°C. Depending on which of the two ions
is present in a higher concentration in a solution, the solution is called to be acidic or alkaline. To express this fact in terms
of a simple number, the negative exponent of the easily measurable hydronium ion concentration [H3O + ] was chosen. This
dimensionless number is called the pH value. The pH can also be described as the negative, decadic logarithm of the hydronium
ion concentration of a solution:
pH = – log [H3O + ]
The hydronium ion concentration of pure water is 10 –7 mol/L, as can be derived from the above equation for the ion product.
Therefore, the pH value is 7.
In an acidic solution, the concentration of H3O + ions is increased (e.g. from 10 –7 mol/L to 10 –2 mol/L) and the pH is therefore
< 7. In an alkaline solution, the hydronium ion concentration is decreased and this results in a pH > 7.
The pKa value
To change the pH value of a solution, substances are dissolved in the solution which release H + ions into the water (acids),
thus raising the H3O + ion concentration and lowering the pH value, or which decrease the H3O + ion concentration and thus
increase the pH value by taking up H + ions (bases). As always in chemistry, this reaction is an equilibrium reaction, and the
capacity of a substance to shift this equilibrium in one direction or another is determined by the potency of the acid. It is
calculated from the equilibrium constants using the following equation
K = [A – ] x [H3O + ] / [HA],
2 Biological Buffers • AppliChem © 2008